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Chapter 10 States of Matter

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Chapter 10 States of Matter

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  1. Chapter 10States of Matter 10.1 The Kinetic-Molecular Theory of Matter

  2. Kinetic-Molecular Theory of Gases • Particles of matter are ALWAYS in motion; constant, rapid motion. (kinetic energy!) • Particles are very small & relatively far apart. • Collisions of particles with container walls cause pressure exerted by gas. • Volume of individual particles is  zero. • Particles exert no attractive or repulsive forces on each other.

  3. Kinetic-Molecular Theory of Gases • Gas particles undergo elastic collisions: Collisions in which no energy is lost • Average kinetic energy is directly proportional to Kelvin temperature of a gas. Air Hockey Table

  4. Ideal Gas • An imaginary gas that perfectly fits all the assumptions of the kinetic-molecular theory • A gas with its particles in constant random motion without attraction for each other is called an Ideal Gas. These particles undergo elastic collisions. • Nearly all real gases behave as ideal gases EXCEPT at very low temperatures or high pressures.

  5. Real Gases • A gas that does not behave completely according to the assumptions of the kinetic-molecular theory. • Real gases occupy space and exert attractive forces on one another

  6. Kinetic-Molecular Theory of the Nature of Gases • Expansion Gases do not have a definite shape or volume Gases take the shape of their containers Gases evenly distribute themselves within a container • Fluidity Gas particles easily flow past one another • Low Density • A substance in the gaseous state has 1/1000 the density of the same substance in the liquid or solid state • Compressibility • Gases can be compressed, decreasing the distance between particles, and decreasing the volume occupied by the gas

  7. Kinetic-Molecular Theory of the Nature of Gases • Diffusion Spontaneous mixing of particles of two substances caused by their random motion • Rate of diffusion is dependent upon: • speed of particles • diameter of particles • attractive forces between the particles

  8. Kinetic-Molecular Theory of the Nature of Gases • Effusion Process by which particles under pressure pass through a tiny opening • Rate of effusion is dependent upon: • speed of particles (small molecules have greater speed than large molecules at the same temperature, so the effuse more rapidly)

  9. Chapter 10States of Matter 10.2 Liquids

  10. Some Properties of a Liquid • Surface Tension: The resistance to an increase in its surface area (polar molecules, liquid metals). A force that tends to pull adjacent parts of a liquid's surface together, thereby decreasing surface area to the smallest possible size.

  11. Some Properties of a Liquid • Capillary Action:Spontaneous rising of a liquid in a narrow tube.

  12. Some Properties of a Liquid • Viscosity: Resistance to flow (molecules with large intermolecular forces).

  13. Some Properties of Liquids Volatility • Liquids that have weak forces of attraction and evaporate easily Nonvolatile Liquids • Liquids that have strong forces of attraction and do not evaporate easily

  14. Properties of Fluids Relative High Density • 10% less dense than solids (average) • Water is an exception • 1000x more dense than gases Relative Incompressibility • The volume of liquids doesn't change appreciably when pressure is applied Ability to Diffuse • Liquids diffuse and mix with other liquids • Rate of diffusion increases with temperature

  15. Chapter 10States of Matter 10.3 Solids

  16. Types of Solids • Crystalline Solids: highly regular arrangement of their components • [table salt (NaCl), pyrite (FeS2)].

  17. Types of Solids • Amorphous solids aka supercooled liquids: considerable disorder in their structures (glass). • Greek for "without shape" • Formation of amorphous solids: • Rapid cooling of molten materials can prevent the formation of crystals * They do not have definite melting points

  18. Representation of Components in a Crystalline Solid Lattice: A 3-dimensional system of points designating the centers of components (atoms, ions, or molecules) that make up the substance.

  19. Types of Crystalline Solids Ionic Solid: contains ions at the points of the lattice that describe the structure of the solid (NaCl).

  20. Unit Cell • The smallest portion of a crystal lattice that shows the three-dimensional pattern of the entire lattice

  21. Types of Crystalline Solids Molecular Solid: discrete covalently bonded molecules at each of its lattice points (sucrose, ice).

  22. Packing in Metals Model: Packing uniform, hard spheres to best use available space. This is called closest packing. Each atom has 12 nearest neighbors.

  23. Closest Packing Holes

  24. Metal Alloys • Substitutional Alloy: some metal atoms replaced by others of similar size. • brass = Cu/Zn

  25. Metal Alloys(continued) • Interstitial Alloy:Interstices (holes) in closest packed metal structure are occupied by small atoms. steel = iron + carbon

  26. Network Solids • Composed of strong directional covalent bondsthat are best viewed as a “giant molecule”. • brittle (non-flexible) • do not conduct heat or electricity • carbon, silicon-based • graphite, diamond, ceramics, glass

  27. Sulfur – S8

  28. Phosphorus – P4

  29. Diamond

  30. Graphite

  31. Zirconia

  32. Chapter 10States of Matter 10.4 Changes of State

  33. Equilibrium • Dynamic condition in which two opposing changes occur at equal rates in a closed system • A closed system at constant temperature will reach an equilibrium position at which the rates of evaporation and condensation will be the same

  34. Equilibrium Vapor Pressure • The pressure of the vapor present at equilibrium. • Determined principally by the size of the intermolecular forces in the liquid. • Increases significantly with temperature. • Volatile liquidshave high vapor pressures. Increasing the temperature will move more particles into the vapor phase to compensate for the new energy

  35. The conversion of a liquid to a vapor within the liquid as well as at its surface. It occurs when the equilibrium vapor pressure of the liquid equals the atmospheric pressure Boiling Boiling Point • The temperature at which the equilibrium vapor pressure of the liquid equals the atmospheric pressure Water boils at 100 °C at 1 atm pressure Water boils above 100 °C at higher pressures Water boils below 100 °C at lower pressures

  36. LeChatelier’s Principle When a system at equilibrium is placed under stress, the system will undergo a change in such a way as to relieve that stress.

  37. Translation: When you take something away from a system at equilibrium, the system shifts in such a way as to replace what you’ve taken away. When you add something to a system at equilibrium, the system shifts in such a way as to use up what you’ve added.

  38. LeChatelier’s Example #1 A closed container of ice and water at equilibrium. The temperature is raised. Ice + Energy  Water The equilibrium of the system shifts to the _______ to use up the added energy. right

  39. LeChatelier’s Example #2 A closed container of N2O4 and NO2 at equilibrium. NO2 is added to the container. N2O4 + Energy  2 NO2 The equilibrium of the system shifts to the _______ to use up the added NO2. left

  40. LeChatelier’s Example #3 A closed container of water and its vapor at equilibrium. Vapor is removed from the system. water + Energy  vapor The equilibrium of the system shifts to the _______ to produce more vapor. right

  41. constant Temperature remains __________ during a phase change. Water phase changes

  42. Phase Diagram • Represents phases as a function of temperature and pressure. • Critical temperature: temperature above which the vapor can not be liquefied. • Critical pressure: pressure required to liquefy AT the critical temperature. • Critical point: critical temperature and pressure (for water, Tc = 374°C and 218 atm).

  43. Water Water

  44. Phase changes by Name

  45. Carbon dioxide Carbon dioxide

  46. Carbon Carbon

  47. Sulfur

  48. Chapter 10States of Matter 10.5 Water

  49. Water’s Properties

  50. Sea Ice • Ice forms on top of the ocean in a thin layer & acts to insulate the warmer waters below from the colder air temperatures. • This occurs in the polar regions, the Artic & Antarctic • Since ice is less dense than liquid water, it will float on top, instead of sinking which would kill all life below the surface. • Sea ice is not the same as an iceberg. Icebergs are pieces of glaciers which are formed by snowfall on land. • Sea ice is not salty, as the hydrogen bonds that hold ice together will not form properly if salt remains in the structure.